CHEMISTRY 59-320CHEMISTRY 59-320ANALYTICAL CHEMISTRYANALYTICAL CHEMISTRY

Fall - 2010Fall - 2010

Lecture 8

6-1 The equilibrium constant

• For an arbitrary reaction

where [A] in the above equation really means [A]/(1 M) if A is a solute, or (pressure in bars)/(1 bar) if A is a gas.

• All terms in the above equation is dimensionless.

• 1. Concentrations of solutes should be expressed as moles per liter.

• 2. Concentrations of gases should be expressed in bars.

• 3. Concentrations of pure solids, pure liquids, and solvents are omitted because they are unity.

• These conventions are arbitrary, but you must use them if you wish to use tabulated values of equilibrium constants, standard reduction potentials, and free energies.

• In the thermodynamic derivation of the equilibrium constant, each quantity in Equation is expressed as the ratio of the concentration of a species to its concentration in its standard state.

• For solutes, the standard state is 1 M. For gases, the standard state is 1 bar (≡105 Pa; 1 atm ≡ 1.013 25 bar).

• For solids and liquids, the standard states are the pure solid or liquid.

The equilibrium constant is more correctly expressed as a ratio of activities.

• If the direction of a reaction is reversed, the new value of K is simply the reciprocal of the original value of K.

• If two reactions are added, the new K is the product of the two individual values.

• If n reaction are added, the overall equilibrium constant is the product of n individual equilibrium constants.

• Examples: Problem 6-6 From the equations HOCl ↔ H+ + OCl- K = 3.0 x 10-8

HOCl + OBr- ↔ HOBr + OCl- K = 15 find the value of K for the reaction HOBr ↔ H+ + OBr-.

All species are aqueous. Answer:

6-2 Equilibrium and thermodynamics

• Enthalpy and entropy affect independently on the degree of a reaction process.

• Gibbs free energy determines whether a reaction process is favoured:

• Reaction Gibbs energy:

• Standard reaction Gibbs energy

0 lnG G RT Q an

activityof productsQ

activityof react ts

00 lnG RT K

• The relationship between equilibrium constant K and thermodynamic properties:

• Le Chatelier’s principle: If a system at equilibrium is subjected to a change that disturbs the system, the direction in which the system proceeds back to equilibrium is such that the change is partially offset.

• Chemical equilibrium constant deals with thermodynamic predictions, not kinetic prediction.

Problem 6-9: The formation of tetrafluoroethylene from its elements is highly Exothermic: 2F2(g) + 2C(s) ↔ F2C=CF2(g)(a) If a mixture of F2, graphite, and C2F4 is at equilibrium in a closed container, will the reaction go to the right or to the left if F2 is added?(b) Rare bacteria from the planet Teflon eat C2F4 and make teflon for their cell wall. will the reaction go to right or to the left if these bacteria are added?(d) Will the reaction go right or left if the container is crushed to one-eighth of its original volume? ( how does the K value change in this case?)(e) Does the equilibrium constant become larger or smaller of the container is heated?Solution: in class discussion

6-3 Solubility product

• The solubility product is the equilibrium constant for the reaction in which a solid salt dissolves to give its constituent ions in solution.

• Solid is omitted from the equilibrium constant because it is in its standard state with an activity value of 1. For example:

• A solution containing excess, undissolved solid is said to be saturated.

• Physical meaning of solubility:• Solubility is governed by more than the solubility product.

Separation by precipitation

• Calculate quotient Q, if Q is larger than equilibrium constant K, precipitation takesplace.

6-4 Complex formation

• A Lewis acid accepts a pair of electrons from a Lewis base when the two form a bond:

• The product of the reaction between a Lewis acid and Lewis base is called an adduct.

• The bond between a Lewis acid and a Lewis based is called a dative or coordinate covalent bond.

• The formation of complex ions increases solubility.

Notation for formation constants

• Formation constants are the equilibrium constants for complex formation.

• The stepwise formation constant, designed Ki, are defined as follows: M + X ↔ MX K1 = [MX]/([M][X]) MX + X ↔ MX2 K2 = [MX2]/([MX][X])

• The overall formation constants are denoted βi, see eqs 6-13 , 6-14, 6-15

βn = K1K2… Kn

6-5 Protic acids and bases

• Acid: a substance that increases the concentration of hydronium ion when added tp water.

• Base: a substance that decreases the concentration of hydronium ion.

HCl + H2O ↔ H3O+ + Cl-

• Bronsted-Lowry acid: proton donor.• Bronsted-Lowry base: proton acceptor HCl(g) + NH3(g) ↔ NH4

+Cl-(s) (acid) (base) (salt)

Conjugate acids and bases

Conjugate acids and bases are related to each other by the gain or loss of one H+.

The nature of H+ and OH-

Autoprotolysis

• Protic solvents have a reactive H+.• All protic solvents undergo autoprotolysis, such

as

•The extent of autoprotolysis is very small.

6-6 pH

6-7 Strengths of acids and bases

Ka=4.2 x 10-13

Ka=6.3 x 10-8

Ka=7.1 x 10-3

Monovalent metal ions are very weak acids (Na+, Ka = 10−13.9). Divalent ions tend to be stronger (Fe2+, Ka = 10−9.4) and trivalent ions are stronger yet (Fe3+, Ka = 10−2.19).